اساسيات التوزيع الإلكتروني - قاعدة البناء التصاعدي ( أفباو ) - أمثلة - أ / طلال الشوافي

Introduction to Electronic Distribution

Overview of the Lesson

• The speaker welcomes students and expresses happiness in connecting through a series of chemistry lessons, specifically focusing on electronic distribution fundamentals.

• Congratulations are extended to high school graduates, and new students for the academic year 2021 are welcomed with wishes for success.

Importance of Electronic Distribution

• The lesson aims to address common issues students face regarding electronic distribution across various educational stages.

• Atoms contain different numbers of electrons, which are arranged around the nucleus in organized fields known as energy levels.

Understanding Energy Levels

Structure of the Atom

• The atom consists of a positively charged nucleus containing protons and neutrons; the charge is determined by protons since neutrons are neutral.

• Electrons are distributed around the nucleus in seven main energy levels, numbered from 1 to 7 and represented by symbols K through Q.

Energy Level Characteristics

• As the principal quantum number (n) increases from 1 to 7, so does the energy associated with that level.

• The first level is closest to the nucleus with lower energy, while the seventh level is farthest away with higher energy.

Sublevels and Orbitals

Division of Energy Levels

• Each main energy level contains sublevels; for example, each principal level has a specific number of sublevels corresponding to its principal quantum number.

• The first main level has one sublevel (1s), while subsequent levels have more:

• Second level: two sublevels (2s, 2p)

• Third level: three sublevels (3s, 3p, 3d)

• Fourth level: four sublevels (4s, 4p, 4d, 4f).

Orbital Capacity

• Each type of orbital can hold a maximum number of electrons:

• s orbitals can hold up to 2 electrons,

• p orbitals can hold up to 6,

• d orbitals can hold up to 10,

• f orbitals can hold up to 14.

Maximum Electron Capacity

Calculating Electron Limits

• The maximum number of electrons in each principal energy level follows the formula 2n^2, where n is the principal quantum number.

• For instance:

• First level holds a maximum of 2 electrons,

• Second holds up to 8 (2 from s + 6 from p),

• Third holds up to 18 (2 from s + 6 from p +10 from d).

This structured approach provides clarity on electronic distribution principles essential for understanding atomic structure in chemistry.

Understanding Electron Configuration

Maximum Number of Electrons in Energy Levels

• The maximum number of electrons that can occupy the fifth energy level is 18, calculated as 2n² where n equals 3.

• This includes eight from the s subshell and ten from the p subshell.

Subshell Structure and Overlap

• The fourth principal energy level (n = 4) contains four subshells: s, p, d, and f.

• Notably, the 4s subshell fills before the 3d due to energy level overlap; this is a critical concept in electron configuration.

Principles Governing Electron Distribution

• Electron distribution refers to how electrons are arranged in atomic orbitals around an atom's nucleus.

• Three main rules govern this distribution: Aufbau principle, Hund's rule, and Pauli exclusion principle.

Aufbau Principle

• The Aufbau principle states that electrons fill orbitals starting from the lowest available energy levels first.

Pauli Exclusion Principle

• According to Pauli's exclusion principle, no two electrons can have identical quantum numbers within an atom.

Visualizing Electron Configuration

• A systematic approach for writing electron configurations involves using a specific order based on increasing energy levels.

• For example, configurations can be expressed using shorthand notation like "1s² 2s²" etc., which simplifies understanding.

Filling Order of Subshells

• The filling order for subshell types begins with s (1 through 7), followed by p (starting at n = 2), then d (starting at n = 3).

Capacity of Subshell Types

• Each subshell has a defined capacity: s holds up to 2 electrons, p holds up to 6, d holds up to 10, and f holds up to 14.

Periodic Table Representation

• The periodic table consists of seven periods representing different principal energy levels.

• Each period ends with a noble gas configuration; for instance:

• Period one ends with Helium (atomic number 2)

• Period two ends with Neon (atomic number 10)

Summary of Energy Levels and Subshell Filling

• Energy levels are numbered from one to seven. Each level corresponds with specific elements in the periodic table.

Practical Application in Electron Configuration

• Understanding these principles allows students to accurately determine electron configurations for various elements based on their position in the periodic table.

Understanding Electron Distribution in Atomic Structure

Overview of Electron Shells and Subshells

• The discussion begins with the introduction of electron shells, specifically focusing on the fourth shell or level (4s).

• The speaker outlines the progression from 3d to 4b, leading into the fifth shell or cycle.

• A detailed explanation is provided about how energy levels are organized from lower to higher, resembling rows in the periodic table from 1 to 7.

Filling Electron Levels

• The filling order starts at the first row and progresses gradually through to the seventh, emphasizing that each subshell fills according to its energy level.

• Transitioning between subshells is highlighted as a key aspect of understanding electron distribution.

• The third subshell is introduced, followed by discussions on filling the fourth (4s), fifth (5s), and their implications for electron distribution.

Aufbau Principle and Noble Gas Configuration

• An explanation of the Aufbau principle illustrates how electrons fill orbitals starting from lowest energy levels before moving to higher ones.

• The use of noble gases as shortcuts for electron configuration is discussed; helium represents the first period while neon marks the end of the second period.

• Argon serves as a reference point for elements beyond atomic number 18, indicating how subsequent shells are filled after noble gases.

Important Rules for Electron Distribution

• Key rules for accurate electronic distribution include ensuring that total electrons match those in a neutral atom and adhering to maximum occupancy limits per orbital.

• Emphasis is placed on maintaining stability during electron distribution, particularly within transition metals where unique configurations may arise due to half-filled or fully filled subshell conditions.

Examples of Electron Configuration

• A practical example using aluminum (atomic number 13) demonstrates detailed step-by-step electronic configuration: 1S^2, 2S^2, 2P^6.

• It’s noted that valence electrons play a crucial role in chemical bonding; thus identifying them post-distribution is essential.

Further Examples and Clarifications

• Another example with calcium (atomic number 20): detailed breakdown shows how it follows similar principles but highlights differences based on atomic structure.

• Bromine's configuration illustrates challenges when selecting noble gas references; careful selection ensures accuracy in representation.

Understanding Electron Configuration and Noble Gases

Key Concepts in Electron Configuration

• The discussion begins with a reference to the number 5, indicating a specific electron configuration or atomic number.

• A mention of noble gases suggests a focus on stable electron configurations, particularly argon as an example.

• The speaker refers to "closest noble gas" which implies using noble gases as a reference point for understanding electron arrangements.

• A series of numbers (18, 20, 30, 35) are presented, likely representing atomic numbers or elements that share similar properties due to their electron configurations.

Elements of the Fourth Period

• Discussion shifts to elements in the fourth period of the periodic table, emphasizing their unique characteristics.

• The seventh group is mentioned, possibly referring to halogens or other groups within the periodic table that exhibit distinct behaviors based on their electron configurations.

Special Cases in Transition Metals

• An exceptional case is introduced regarding transition metals where electronic distribution can lead to unusual stability; specifically mentioning chromium's configuration (4s1 3d5).

• Another special case highlights how certain transition metals may have electrons shift from s-orbitals to d-orbitals for increased stability; this includes adjustments in their electronic distribution for optimal stability.